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Chemistry: Atoms First
Second Edition
Julia Burdge & Jason Overby
Chapter 4
Periodic Trends
of the Elements
M. Stacey Thomson
Pasco-Hernando State College
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
4.1
Development of the Periodic Table
In 1864, John Newlands noted that when the elements were
arranged in order of atomic number that every eighth element had
similar properties.
 He referred to this as the law of octaves.
In 1869, Dmitri Mendeleev and
Lothar Meyer independently
proposed the idea of periodicity.
Mendeleev grouped elements
(66) according to properties.
Mendeleev predicted properties
for elements not yet discovered,
such as Ga.
4.2
Classification of Elements
The main group elements (also called the representative elements)
are the elements in Groups 1A through 7A.
Classification of Elements
The noble gases are found in Group 8A and have completely filled
p subshells.
The Modern Periodic Table
The transition metals are found in Group 1B and 3B through 8B.
 Group 2B have filled d subshells and are not transition metals.
The Modern Periodic Table
The lanthanides and actinides make up the f-block transition
elements.
The Modern Periodic Table
There is a distinct pattern to the electron configurations of the
elements in a particular group.
The Modern Periodic Table
The electrons in the outermost occupied PRINCIPLE ENERY
LEVEL of an atom are called the valence electrons.
Valence electrons are involved in the formation of chemical bonds.
For Group 1A: [noble gas]ns1
core
For Group 2A: [noble
valence
gas]ns2
core
valence
For Group 7A: [noble gas]ns2np5
core
valence
Example: [He]2s1
core valence
Example: [Ar]4s2
core valence
Example: [Ne]3s23p5
core valence
Worked Example 4.2
Without using a periodic table, give the ground-state electron configuration and
block designation (s-, p-, d-, or f-block) of an atom with (a) 17 electrons, (b) 37
electrons, and (c) 22 electrons. Classify each atom as a main group element or
transition metal.
4.3
Effective Nuclear Charge
Effective nuclear charge (Zeff) is the actual magnitude of positive
charge that is “experienced” by an electron in the atom.
In a multi-electron atom, electrons are simultaneously attracted to
the nucleus and repelled by one another.
 This results in shielding, where an electron is partially shielded
from the positive charge of the nucleus by the other electrons.
 Although all electrons shield one another to some extent, the
core electrons shield the most.
 As a result, the value of Zeff increases steadily from left to right
because the core electrons remain the same but Z increases.
Z
Zeff
Li
Be
B
C
N
O
F
3
4
5
6
7
8
9
1.28
1.91
2.42
3.14
3.83
4.45
5.10
Atomic radii (in picometers)
4.4
Atomic Radius
The Atomic radius: the
distance from the atom’s
nucleus and its valence
shell.
Atomic radius increases
from top to bottom because
outermost shell lies farther
from the nucleus
Atomic radius decreases
from left to right because
of increasing Zeff which
draws the valence shell
closer to the nucleus
Worked Example 4.3
Referring only to a periodic table, arrange the elements P, S, and O in order of
increasing atomic radius.
Strategy Use effective nuclear charge to compare the atomic radii of two of the
three elements at a time.
Text Practice: 4.14 4.20 4.30 4.44
Ionization Energy
Ionization energy (IE) is the minimum energy required to remove
an electron from an atom in the gas phase.
The result is an ion, a chemical species with a net charge.
Na(g) → Na+(g) + e−
Sodium has an ionization energy of 495.8 kJ/mol.
Specifically, 495.8 kJ/mol is the first ionization energy of sodium,
IE1(Na), which corresponds to the removal of the most loosely
held electron.
Ionization Energy
Ionization Energy
In general, as Zeff increases, ionization energy also increases.
 Thus, IE1 increases from left to right across a period.
Ionization Energy
Within a given shell, electrons with a higher value of l are higher
in energy and thus, easier to remove.
Ionization Energy
Removing a paired electron is easier because of the repulsive
forces between two electrons in the same orbital.
Ionization Energy
It is possible to remove additional electrons in subsequent
ionizations, giving IE1, IE2, and so on.
Na(g) → Na+(g) + e−
IE1(Na) = 496 kJ/mol
Na+(g) → Na2+(g) + e−
IE2(Na) = 4562 kJ/mol
Ionization Energy
It takes more energy to remove the 2nd, 3rd, 4th, etc. electrons
because it is harder to remove an electron from a cation than an atom.
It takes much more energy to remove core electrons than valence.
 Core electrons are closer to nucleus.
 Core electrons experience greater Zeff because of fewer filled
shells shielding them from the nucleus.
Worked Example 4.4
Would you expect Na or Mg to have the greater first ionization energy (IE1)?
Which should have the greater second ionization energy (IE2)?
Strategy Consider effective nuclear charge and electron configuration to
compare the ionization energies. Na has one valence electron and Mg has two.
Text Practice: 4.50
Electron Affinity
Electron affinity (EA) is the energy released when an atom in the
gas phase accepts an electron.
Cl(g) + e− → Cl−(g)
Electron Affinity
Like ionization energy, electron affinity increases from left to right
across a period as Zeff increases.
 Easier to add an electron as the positive charge of the nucleus
increases.
Electron Affinity
It is easier to add an electron to an s orbital than to add one to a p
orbital with the same principal quantum number.
Electron Affinity
Within a p subshell, it is easier to add an electron to an empty
orbital than to add one to an orbital that already contains an
electron.
Worked Example 4.5
For each pair of elements, indicate which one you would expect to have the
greater first electron affinity, EA1: (a) Al or Si.
Strategy Consider effective nuclear charge and electron configuration to
compare the ionization energies. (a) Al is in Group 3A and Si is in Group 4A. Al
has three valence electrons ([Ne]3s23p1), and Si has four valence electrons
([Ne]3s23p2).
Text Practice: 4.58
Study Guide for Sections 4.1-4.4
DAY 9, Terms to know:
Sections 4.1-4.4 valence electrons, core electrons, effective nuclear charge (Zeff),
shielding, atomic radii, ionization energy, electron affinity
DAY 9, Specific outcomes and skills that may be tested on exam 1:
Sections 4.1-4.4
•Be able to give a complete or abbreviated electron configuration for an atom in
either its ground state or one possible excited state
•Be able to give a complete or abbreviated orbital diagram for an atom either its
ground state or one possible excited state
•Be able to describe what effective nuclear charge is and how it is calculated
•Be able to rank relative atomic radii, electron affinity, ionization energy ionic radii,
and explain WHY they are ranked based on attractions and repulsions within the
atom
•Be able to rank atoms in order of greatest ionization energies including IE1, IE2,
IE3, etc. and explain WHY they should be ranked in that order
Extra Practice Problems for Sections 4.1-4.4
Complete these problems outside of class until you are confident you have learned
the SKILLS in this section outlined on the study guide and we will review some of
them next class period. 4.17 4.19 4.45 4.49 4.51 4.55 4.57 4.59 4.61 4.63
4.107 4.127 4.129
Prep for Day 10
Must Watch videos:
https://www.youtube.com/watch?v=Qf07-8Jhhpc (Tyler DeWitt: ionic bonds part 1)
https://www.youtube.com/watch?v=5EwmedLuRmw (Tyler DeWitt: ionic bonds part 2)
https://www.youtube.com/watch?v=RkZNYuSho0M (Tyler DeWitt: ionic bonds part 3)
https://www.youtube.com/watch?v=X_LVANMpJ0c (Tyler DeWitt: writing ionic formulas)
Other helpful videos:
https://www.youtube.com/watch?v=6GjYGdk32U&list=PLqOZ6FD_RQ7kTjN4O2MNzf5YfeiIx7SGI (UC-Irvine watch first 15 minutes)
https://www.youtube.com/watch?v=dxSMy4CIwGQ&list=PLqOZ6FD_RQ7k3kp5B4jQbIA99gh73RF
sh (UC-Irvine ions and molecules)
Read Sections 4.5-4.6, 5.1-5.4